17 May 2023, 101383
Article
Synthesis of stable single-crystalline carbon dioxide clathrate powder by pressure swing crystallization
Author links open overlay panelZhiling Xiang 1 5, Congyan Liu 1 5, Chunhui Chen 1, Xin Xiao 2, Thien S. Nguyen 3, Cafer T. Yavuz 3, Qiang Xu 2 4, Bo Liu 1 6
Summary
Reversible CO2 capture and release under ambient conditions is crucial for energy-efficient carbon capture and storage. Here, we report the pressure swing crystallization of CO2 in a single-crystalline guanidinium sulfate-based clathrate salt under practical conditions of 52 kPa and 298 K, with a high CO2 density (0.252 g cm−3) and capacity (17 wt %). The captured CO2 is released as a pure stream through moderate means of pressure or temperature stimulation, all while the desorbed Gua2SO4 is ready for another cycle. The clathrate is selective exclusively to CO2 even in the presence of common flue gas components, such as water vapor and N2, owing to the specific electrostatic interaction between the CO2 and guanidinium cations. The mechanism unraveled through single-crystal studies is distinctively different from physisorption or chemisorption, opening up a promising venue for future carbon capture and storage technologies through rapid CO2 solidification using an abundant salt.
Graphical abstract
Keywords
carbon captureadsorptionCO2 storageguanidinium sulfategas hydratessustainable chemistryflue gas treatmentdirect air capture
Introduction
Reversible carbon capture and release in an energy-efficient way is vital toward many industrial processes, particularly in efforts to address the current global warming crisis. In general, CO2 capture solutions can be classified into physisorption and chemisorption processes through temperature or pressure swing operations. But each strategy possesses its own merits and shortcomings. Physisorption of CO2 via weak interactions mainly involves porous materials with high surface areas, such as porous carbons,1,2 zeolites,3,4 metal-organic frameworks,5,6,7 covalent-organic frameworks,8 and hydrogen-bonded organic frameworks,9 giving rise to low sorption heat and easy adsorbent regeneration. Parasitic molecules like water, however, compete with CO2 and deteriorate the selectivity, capacity, and cycling performance for carbon capture and ultimately increase the energy consumption for regeneration in practical scenarios.10,11,12,13 Chemisorption of CO2, on the other hand, usually produces high heat and therefore requires intensive energy input for absorbent regeneration.14,15,16 This is because the captured CO2 is converted into carbamates and carbonates (HCO3− or CO32−) when brought into contact with aqueous amine solutions or tethered amines on porous supports. One example is guanidine, a multi-amine construct, which has been explored as a chemisorptive CO2 absorbent for direct carbon capture from air.17,18,19,20 In nature, guanidine-based CO2 capture is a chemical conversion process taking advantage of basicity of guanidine, similar to carbon capture using alkaline solutions or monoethanolamine. This chemical product required high regeneration energy, usually at temperatures higher than 120°C. An interesting example is a 3D hydrogen-bonded framework assembled from tetrahedral tetraamidinium cations and carbonates via nonelectrostatic hydrogen bond in water,21 but utility for a cyclic CO2 capture through dynamic complexation is not clear.
Gas hydrates such as CO2 hydrates are suitable for a rapid, reversible carbon capture. They are, however, often generated at low temperatures and high pressures (for example, T = 0°C and P = 1,200 kPa). In a CO2 hydrate, CO2 is enclosed in a water cage constructed through hydrogen bonds.22 In other words, CO2 forces icy water to crystallize into a hydrate framework, where it is also known to trap other guest species such as CH4, N2, or small organic molecules.23 Under increasing temperature or reduced pressure, for example bringing CO2 clathrates to ambient conditions, the water cages collapse, and trapped guest molecules escape. Laboratory syntheses of gas hydrates have been accomplished.24 The kinetics of clathrate formation, however, remain unchanged, where the conditions of low temperature and high pressure are indispensable, making CO2 hydrates impractical for applications of gas adsorption, separation, and storage. Recently, we reported a reversible structural transformation of [B(OCH3)4]3[C(NH2)3]4Cl⋅4CH3OH upon MeOH capture and release as an example of achieving dynamic behavior of gas hydrates at ambient conditions.25 Carbon capture and storage as a powder based on clathrate formation using CO2 hydrate is challenging and remains elusive, to the best of our knowledge.
Herein, we report a stable CO2 clathrate powder formation under ambient conditions through co-crystallization of CO2 with guanidinium sulfate (Gua2SO4, where Gua is guanidinium) from an aqueous Gua2SO4 solution (Figure 1). As revealed by our single-crystal studies, crystallization of CO2@Gua2SO4 is triggered by dominant electrostatic interactions between CO2 and guanidinium ions, which are encased among the strong hydrogen bond interactions between guanidinium cations and sulfate. CO2@Gua2SO4 readily releases CO2 through structure collapsing in ambient conditions, and the resultant Gua2SO4 is ready for another cycle without requiring further regeneration (Figure 1). The volume and weight densities of CO2 in CO2@Gua2SO4 are determined to be 0.252 g cm−3 and 17 wt %, respectively, revealing its tremendous potential for carbon capture and storage in practical conditions.
Figure 1. Schematic demonstration of reversible clathrate formations with pressure swing for CO2 capture and release process
(A) CO2 hydrate where CO2 molecules are trapped in water clusters at high pressures and low temperatures.
(B) A clathrate from guanidinium sulfate and CO2, CO2@Gua2SO4, where CO2 can be captured at pressures as low as 32 kPa and temperatures at flue gas conditions (35°C and below).
Results
Synthesis and crystal structure of CO2@Gua2SO4
Upon charging an aqueous Gua2SO4 solution with CO2, we observe spontaneous formation of a single-crystalline CO2@Gua2SO4 (see experimental procedures for details; Figure S1–S7, 2A, and 2B). The crystal structure of CO2@Gua2SO4 is determined by single-crystal X-ray diffraction experiments at 100 K (Figure S2; Table S1). In CO2@Gua2SO4, each guanidinium ion adopts two sets of hydrogen bonds (H-bonds) connecting three SO42− ions (Figure 2C), while each sulfate ion connects six guanidinium ions via multiple H-bonds (Figure 2F). The extended H-bond system (Figure S3; Table S2) in 3D results in a H-bonded framework crystallized at a tetragonal space group in which four CO2 molecules are accommodated in each unit cell (Figures 2E, 2F, and S4). As shown in Figure S5, an irregular cage comprised of five sulfate and eight guanidinium ions can be identified, in which one CO2 molecule is accommodated, and the cages are stacked by plane sharing. In contrast, free Gua2SO4 crystallizes with a cubic space group in a dense-stacking mode, in which the H-bond connection modes between Gua+ and SO42− ions are notably different (Figure S7).
Figure 2. Formation and single-crystal structure of the clathrate, CO2@Gua2SO4
(A) Photos of CO2@Gua2SO4 precipitation under CO2 atmosphere and magnetic stirring. Test tube diameter is 2.8 cm.
(B) Optical image of CO2@Gua2SO4 crystal. Scale bar is 500 μm.
(C) Hydrogen-bond interactions of guanidinium cations with three SO42− ions.
(D) Hydrogen-bond interactions of SO42− with six guanidinium ions.
(E and F) 3D hydrogen-bonded framework of CO2@Gua2SO4 omitting CO2 (E) and (F) with captured CO2. Dash lines represent hydrogen bonds. Hydrogen-bond interaction is deduced from single-crystal X-ray diffraction measurements. Hydrogens are omitted for clarity. Color code: gray, carbon; blue, nitrogen; red, oxygen; yellow, sulfur.
Surprisingly enough, in CO2@Gua2SO4, a H-bond interaction between CO2 and Gua2SO4 is not favored because the distances and angles are out of range for a typical H-bond interaction (Figure S8), and this creates a tremendous opportunity for a reversible CO2 capture. In a molecule of CO2, C and O atoms bear partially positive and negative charges, respectively. The C atom exists as a carbocation and N atoms bear partial negative charge in a guanidinium cation. The distance among these oppositely charged atoms ranges from 3.4 to 3.9 Å (Figure 3A) so that the electrostatic interaction is more pronounced. One CO2 molecule interacts with three guanidinium ions with a “triple-team structure,” ensuring an effective coulombic force. The distance between positive C in CO2 and negative O in SO42− is determined to be 5.026 Å (Figure 3B), indicating a much weaker electrostatic interaction. We ascribe these multiple electrostatic interactions as the main driving forces for CO2-induced crystallization of CO2@Gua2SO4 from aqueous solution under CO2 atmosphere. The solid-state magic-angle spin 13C nuclear magnetic resonance (NMR) spectrum of CO2@Gua2SO4 displays a sharp chemical shift at 124.9 ppm (Figure 3C), equal to the chemical shift of CO2 in a physisorption state.26,27 This is consistent with the result from crystal structure analysis as mentioned above. The other peak at 159.2 ppm is assigned to the carbocation in the guanidinium. A similar peak with stretching vibration of gas phase CO2 appears in the infrared (IR) spectrum of CO2@Gua2SO4 at 2,335 cm−1, which also suggests the weak interaction of CO2 with guanidinium and SO42− ions (Figure 3D). It is worth noting that when reacting with hydroxyl and amine functional groups, CO2 is often converted into carbonate ester or carbamate.14,15,16,28 Owing to the positive charge, guanidinium cations with three amine groups interact with CO2 via an electrostatic interaction instead of a chemical reaction. This moderate but abundant interaction is strong enough to pull guanidiniums and SO42− together for CO2 clathrate formation and crystallization in aqueous solution. In synthetic terms, the CO2 inclusion mechanism unraveled by single-crystal structure analysis manifests that amino groups could interact with CO2 via electrostatic interactions instead of a chemical reaction, owing to the delocalization of lone-pair electrons from amino functionalities to carbocations of guanidinium ions. The results point to a substantial parameter space to tune the interactions between CO2 and sorbents for optimized carbon capture in terms of energetics and economics.
Figure 3. Interaction of CO2 with guanidinium and sulfate ions
(A) Electrostatic interactions between CO2 and guanidinium cations.
(B) Distance between CO2 and sulfate ions in CO2@Gua2SO4 based on single-crystal X-ray diffraction data.
(C) Solid-state magic-angle spin 13C NMR spectrum of CO2@Gua2SO4 and Gua2SO4.
(D) ATR-IR spectra of CO2@Gua2SO4 and Gua2SO4.
In order to screen counter anions on guanidinium, we have tested a series of guanidinium salts for CO2 clathrate potential, including nitrate, chloride, and phosphate (see details in experimental procedures). None, however, resulted in any precipitate under the same conditions using guanidinium cations. Chloride ion is not favorable for H-bond formation,29 while nitrate ion and guanidinium generate a plane structure, as both ions are planar,30 and therefore not beneficial for CO2 clathrate formation. The complicated hydrolysis process of phosphate in aqueous solution may account for the failure of CO2 co-precipitation.31 It is safe to conclude that the complex H-bond interactions between guanidinium and sulfate are crucial for crystal formation and CO2 capture.
Optimization of CO2 clathrate formation
In an attempt to explore the boundary conditions of CO2@Gua2SO4 formation, we varied pH, Gua2SO4 concentration, CO2 pressure, and temperature. Since basic aqueous solutions directly react with CO2, we tuned the pH of Gua2SO4 solutions from 1 to 7 using a H2SO4 solution. We found that pH value has little influence over CO2@Gua2SO4 formation (Figure 4A). Considering the possible influence of anion size on boundary conditions, HCl and HNO3 were also used to adjust the pH. The results showed that the sizes of the anions exert negligible effects on the formation conditions of CO2@Gua2SO4 (Figure S9). Increasing CO2 pressure led to a higher production rate of CO2@Gua2SO4 (Figure S10). The saturated Gua2SO4 concentration is determined to be 72 and 75.7 wt % in water at 273 and 298 K, respectively. In order to eliminate the influence of the Gua2SO4 concentration variation during the continuous precipitation of CO2@Gua2SO4, we used a saturated solution in equilibrium with excess Gua2SO4 powder to investigate the temperature-pressure relationship under phase equilibrium conditions. At equilibrium, the CO2 pressure-temperature relationship fitted well to the Clapeyron-Clausius equation (Figures 4B, S11, and S12), revealing the boundary conditions of CO2@Gua2SO4 formation and indicating a wide range of conditions for capturing CO2 using aqueous Gua2SO4. CO2@Gua2SO4 started to crystallize at a CO2 pressure of 32 and 52 kPa at 273 and 298 K, respectively (Figure 4B), whereas conventional CO2 hydrates are formed under 1,500 kPa at 275 K, about 50 times higher than that of CO2@Gua2SO4 formation.32 The sorption enthalpy change for CO2@Gua2SO4 is calculated to be 15.37 kJ/mol (Figure S12). This value is considerably lower than the enthalpy changes of most CO2 absorbents as summarized in Figure S1319,33,34,35,36,37 and falls into the physisorptive domain (<40 kJ/mol), suggesting minute heat requirements during CO2 capture and release and corresponding to the ease of clathrate formation and collapse. In contrast, the enthalpy change for conventional CO2 hydrate decomposition is 57.1 kJ/mol, owing to the necessity of breaking ample H-bonds among H2O molecules,32 about 3.7 times higher than that of CO2@Gua2SO4.
Figure 4. CO2 sorption behavior in a Gua2SO4 aqueous solution
(A) pH effect on CO2@Gua2SO4 formation.
(B) Pressure-temperature correlation of CO2-Gua2SO4 aqueous system under equilibrium. CO2 hydrate data is plotted as a reference.29
(C) CO2 sorption profile in aqueous Gua2SO4 with or without stirring. CO2 pressure changes of the control experiments, (i) CO2 sorption in pure water with stirring, (ii) N2 sorption in pure water, and (iii) aqueous Gua2SO4 with or without stirring, are given as a reference.
(D) Cyclic performance of CO2 sorption in aqueous Gua2SO4. The margin of error points at a ±3.5% variation.
(E) The pressure changes vs. time plot of CO2 adsorption and desorption in water and aqueous Gua2SO4 with a temperature swing between 273 (adsorption) and 303 K (desorption). Dashed lines are for pure water, and solid lines represent aqueous Gua2SO4.
(F) PXRD data for structural evolution from CO2@Gua2SO4 to Gua2SO4 in air. PXRD patterns of CO2@Gua2SO4 simulated from crystallographic data and as-prepared Gua2SO4 are given for reference (peaks marked as blue and orange belong to CO2@Gua2SO4 and Gua2SO4, respectively).
The CO2 in CO2@Gua2SO4 solidifies at near-ambient conditions with four CO2 molecules accommodated in each unit cell with a cell volume of 1,159 Å3, corresponding to a CO2 volume density of 0.252 g cm−3. This is equal to about 140 m3 CO2 in one cubic meter of CO2@Gua2SO4. Gravimetrically, the CO2 loading in CO2@Gua2SO4 is calculated to be 17 wt %. As shown in Figure S14, 2.3 g CO2 was stored in a 20 mL glass bottle, which corresponds to 1.2 L CO2 gas at standard conditions at room temperature. In contrast, to store the same amount of CO2 gas in a 20 mL bottle, the pressure will have to reach 6,000 kPa at 0°C. Hence, CO2@Gua2SO4 is an ideal candidate for CO2 storage and transport. Other reported CO2 clathrates, in general, are formed under low-temperature and/or high-pressure conditions, as summarized in Table S3. The inability to operate in ambient conditions renders these clathrates not feasible for flue gas treatment, as large deviation of temperature and/or pressure from ambient conditions means substantial energy input for a practical carbon capture operation. Gua2SO4 is perfectly optimized for such a job, as it stores 140 m3 CO2 in 1 m3 clathrate with a gravimetric capacity of 17% (uptake at 52 kPa and room temperature [RT] and release at ambient conditions), positioning it even better than most common porous materials as summarized in Table S4. And the structure does this in a simple pressure swing adsorption (PSA) cycle, eliminating energy-intensive processes and showing quantitative capture of CO2 in the presence of N2 and H2O (as discussed in the next section). It is important to note that only liquid amine solutions (industrial standard) could do such a performance with the kinetics that a process needs.
Kinetics of CO2 clathrate formation
In an isochoric experiment, after removing air by CO2 flushing, the stainless-steel autoclave containing Gua2SO4 solution (1.5 g/g, 60 wt %) was charged with CO2 to 1,000 kPa at 273 K with constant stirring (Figure S15). The pressure drop was monitored against time, and we found that the pressure decreases in an approximatively linear fashion until 60 min and gets stable at about 100 min (Figure 4C). Replacing Gua2SO4 solution with pure water leads to a much smaller pressure drop. The latter is about one-fifth of the aqueous Gua2SO4 solution under the same conditions and results from CO2 dissolution in water and the temperature drop from RT to 273 K. N2, however, does not dissolve well and shows similar behavior in both pure water and Gua2SO4 solutions, proving that N2 cannot form any clathrates under the same conditions. Therefore, we concluded that the structure alternate between Gua2SO4 and CO2@Gua2SO4 exclusively for CO2, even in the presence of N2 and water. The final conversion yield of CO2@Gua2SO4 is around 85 wt % of Gua2SO4 according to the weight increase of the solution.
The kinetics of CO2@Gua2SO4 formation closely follows the Gua2SO4 concentration, CO2 pressure, temperature, and stirring. In principle, the formation of CO2@Gua2SO4 through the steps of gas dissolution, diffusion, nucleation, and growth is quite similar to a conventional CO2 hydrate formation process.38 In a static experiment without stirring, the CO2 pressure drop is very low and similar to the case of gaseous CO2 dissolution in water under the same period (Figure 4C). A crystalline layer was observed over the solution, preventing further gas diffusion. We recorded the crystallization process by sustaining a CO2 pressure at 300 kPa while constantly stirring (Video S1). The solution is initially clear, with CO2 dissolving and diffusing. We note that the nucleation first occurs near the vortex of stirring, which correlates to high CO2 concentrations in the vicinity. Once the nucleation is initiated, the solution becomes turbid very quickly, owing to the rapid crystal growth. Therefore, it is safe to conclude that CO2 dissolution and diffusion are the rate-determining steps for CO2@Gua2SO4 formation.
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Video S1. CO2@Gua2SO4 formation under 0.3 MPa (3 bar) CO2 pressure. The movie is set to play at an increased speed of 20 times
Even though physisorptive CO2 capture processes are well known to be durable with long cycle life, we repeated the sorption process for ten cycles, and the data exhibit excellent reproducibility, as Gua2SO4 is stable and nonvolatile and does not degrade or become lost during the continuous operations (Figures 4D, S16, and S17A). In a temperature swing experiment, we warmed up the autoclave to 303 K after adsorption reached equilibrium at 273 K and observed the pressure increasing (Figure 4E). This is clearly associated with CO2@Gua2SO4 collapse and CO2 release. The final pressure is slightly higher than initial value owing to the higher final temperature of the autoclave. The results reveal a reversible formation and decomposition of a salt-CO2 clathrate accompanied by CO2 capture and release near RT.
Cyclic CO2 capture
The precipitates of CO2@Gua2SO4 can be separated via centrifugation or a simple filtration. In open air, dry CO2@Gua2SO4 decomposes with CO2 release, as verified by powder X-ray diffraction (PXRD), but the decomposition in solid state is sluggish (Figure 4F). The PXRD of the fresh CO2@Gua2SO4 sample matched well with the pattern simulated from single-crystal (SC) XRD data, with some deviation at 15.0° owing to its deposition method at sample preparation and the difference of measurement temperature (Figure S6). Note that the main peaks at 14.6° and 15.3° of CO2@Gua2SO4 gradually fade, while new peaks at low angles of 11.2°, 12.3°, and 15.0° belonging to the Gua2SO4 phase appear and grow with prolonged exposure time in air. The thermogravimetric analysis (TGA) and differential scanning calorimetry (DSC) analyses of the CO2@Gua2SO4 sample in air are shown in Figure S17B. The weight loss of ca. 13.7% before 86.5°C is originated from the removal of CO2, which is slightly less than 17% (theoretical capacity of CO2 in CO2@Gua2SO4). It can be attributed to CO2@Gua2SO4 slightly decomposing with CO2 release during sample preparation and measurements. Upon adding water to a CO2@Gua2SO4 powder, we observe rapid CO2 release with vigorous bubbling, and the solids disappear because Gua2SO4 salt, the only decomposition product, is highly soluble (Video S2). Judging from the shrinking solids, we project the CO2 release proceeds from the outside in. The formed Gua2SO4 on the surface of CO2@Gua2SO4 could block further CO2 release and slow down the decomposition rate, but as Gua2SO4 is dissolved out with water, the CO2 release remains unaltered. The decarbonized Gua2SO4 solid or solution is ready for reuse without requiring any further regeneration step. The released CO2 purity from CO2@Gua2SO4 is 100% except when regeneration was done by water and therefore saturated with water vapor and can be used directly in the CO2 market including in refrigerant, beverage, pharmaceutical, and food storage sectors.
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Video S2. Accelerated CO2@Gua2SO4 collapse and CO2 release with water addition.
CO2 removal from flue gas
Since CO2 capture is conducted in an aqueous solution and moisture exerts no influence over CO2 capture, we tested CO2 removal performance of a Gua2SO4-saturated solution with a dry simulated flue gas of different molar ratios of N2 and CO2. The CO2 percentage in a flue gas is typically in the range of 5%–15% depending on the industrial processes and sources.39 With limited CO2 partial pressure, we pressurized flue gas in order to reach a high CO2 presence. A schematic illustration of setup is shown in Figure 5A. Figure 5B and Table S5 summarize CO2 capture performances at varying total flue gas pressures. Obviously, the higher pressure leads to higher CO2 removal. In the case of a CO2 percentage at 15%, total flue gas needs to be pressurized to 5,000–6,000 kPa to meet DOE targets in terms of a 90% CO2 capture. When the CO2 percentage is increased to 30% (i.e., from the cement industry), the flue gas pressure can be as low as 2,000–3,000 kPa for a 90% CO2 removal. We also note that the CO2 content in natural gas streams can reach over 40% in some cases.40 The high CO2 content will lower the total pressure of mixtures and therefore reduce the CO2 removal cost based on the physisorptive clathrate mechanism of the CO2@Gua2SO4.
Figure 5. CO2 removal performance of Gua2SO4-saturated solution with a dry simulated flue gas
(A) Schematic illustration of the setup.
(B) CO2 removal percentages against total flue gas pressure with CO2 mole percentages at 15 and 30 mol %.
Discussion
In this work, we reported the first ambient example of mimicking CO2 hydrate structure with CO2 as guest molecules. A simple Gua2SO4 co-crystallizes with CO2 into a stable clathrate (CO2@Gua2SO4). The clathrate releases CO2 on demand, and both adsorption and desorption can occur at mild conditions. The H-bonded framework assembled from guanidinium and sulfate mimics the water cage in a CO2 hydrate, effectively trapping four CO2 molecules per unit cell of the clathrate. Electrostatic interactions between guanidinium ion and CO2 are revealed through a single crystal (SC) study instead of the van der Waals interactions commonly observed for CO2 and water cages of CO2 hydrates.
The in situ process of co-crystallization introduced here is fundamentally different from CO2 adsorption in a preformed H-bonded organic framework and is quite unlike CO2 physisorption in other porous materials. A CO2 clathrate imitating the CO2 hydrate, therefore, exhibits unique advantages toward carbon capture. First, Gua2SO4 exclusively captures CO2 without water or moisture interference, swiftly overcoming the fatal weakness of any physisorption processes. Second, CO2 release through structure collapsing could be triggered at ambient conditions, requiring little energy input for absorbent regeneration while defeating the parasitic energy dilemma of chemisorption. Third, Gua2SO4 is stable and noncorrosive, a highly desirable feature when compared with ethanol amine, ammonia, and other basic solutions that are commonly used in carbon capture. Fourth, a stable CO2@Gua2SO4 in powder form is also beneficial for storage and transportation of CO2, benefiting from its remarkably high volume per weight capacity. In addition, by the lessons learned in this study, we can tune dynamic H-bonded frameworks with enriched structural variation, enabling us further to regulate and control the properties for further improving CO2 capture in terms of stability, recyclability, sorption capacity, and selectivity, while lowering regeneration energy penalty and cost.
Experimental procedures
Resource availability
Lead contact
Further information and request for the resources are available from the lead contact, Bo Liu (liuchem@ustc.edu.cn).
Materials availability
No unique materials were generated by this study.
Materials
All chemicals and reagents were purchased from commercial suppliers and used without further purification. Deionized water was used as solvent. Ethanol and sulfuric acid were purchased from Sinopharm Chemical Reagent. Guanidinium carbonate (Gua2CO3) was purchased from Shanghai Adamas Reagent. Nitrogen gas (99.999%), carbon dioxide (99.9%), and carbon dioxide-nitrogen mixture gas (15 mol % CO2+85 mol % N2 and 30 mol % CO2+70 mol % N2) were purchased from Nanjing Special Gas Plant.
Characterization
PXRD measurements were conducted on a Rigaku MiniFlex 600 diffractometer using Cu Kα radiation (λ = 1.5418 Å). The FTIR spectrum was measured by a Nicolet iS5 spectrophotometer with an attenuated total reflectance (ATR) module (Thermo Fisher Scientific, Waltham, MA, USA). Solid-state NMR spectra were recorded by a Bruker AVANCE NEO 600WB spectrometer with a magic-angle spin (MAS) rate of 15 kHz. The 13C resonance frequency was 150.93 MHz at the experiment. 13C single-pulse spectrum was recorded with 16 s pulse delay. TGAs were performed from 25°C to 800°C at a heating rate of 10°C/min in air on a TGA Q500 integration thermal analyzer (For CO2@Gua2SO4, the heating rate from 20°C to 300°C was 5°C/min and from 300°C to 700°C, 10°C/min). Gas chromatography (GC) was performed on a Shimadzu GC 2014 gas chromatograph fitted with a Porapak Q column and TCD detector. Ar gas (99.999%) was used as the carrier gas.
Synthesis of Gua2SO4
Gua2SO4 was synthesized simply by neutralizing guanidinium carbonate with H2SO4. Gua2CO3 (180 g, 1 mol) and deionized water (250 mL) were first mixed in a 1,000 mL beaker. Concentrated H2SO4 (98%, 54 mL, 1 mol) was then added dropwise into the beaker at a rate of one drop per second. The final pH was adjusted to 7 using only H2SO4. Absolute ethanol was added as nonsolvent under constant stirring to precipitate the product. The colorless powder was collected by vacuum filtration, followed by washing with absolute ethanol and drying in an oven at 100°C before use. Yield was calculated to be 99% based on guanidinium. The production of Gua2SO4 was verified by common characterization tools, particularly PXRD (Figure S1).
Synthesis of CO2@Gua2SO4 SC
A CO2@Gua2SO4 SC was grown by treating an aqueous Gua2SO4 solution with CO2 at low temperature. Gua2SO4 (3 g) was dissolved in water (5 mL). After complete dissolution, the solution was added in a glass vial, which was placed in an autoclave. The autoclave was rinsed using CO2 to completely replace the air, and the final CO2 pressure in the autoclave was fixed at 700 kPa. The autoclave was kept at 2°C by a cooling jacket. Colorless plate-like SCs at centimeter sizes were obtained after 12 h.
SC XRD measurements
Since the crystals of CO2@Gua2SO4 readily decompose at atmospheric conditions, a fresh crystal from the reactor was placed on top of a glass fiber under a stream of liquid nitrogen and rapidly mounted onto a sample goniometer for centering. SC XRD data were collected by a Rigaku Oxford Diffraction Super-Nova diffractometer using Mo-Kα radiation (λ = 0.71073 Å) at 100 K. The data collection and processing were carried out with CrysAlisPro software. The crystal structure was solved by direct method and refined by full-matrix least squares based on F2 using an SHELXTL 14XL program package. Hydrogen atoms were fixed geometrically at their positions and allowed to ride on parent atoms. Crystallographic and structure refinements data for CO2@Gua2SO4 are summarized in Table S1.
PXRD of CO2@Gua2SO4
RT PXRD spectra of CO2@Gua2SO4 slightly deviated from the simulated one especially at (004) plane (Figures S5 and S6) because the measurements were conducted at different temperatures. At 100 K, the guest CO2 molecules stopped moving and therefore were forced to align with guanidinium to maximize the H-bonding interactions. At RT, however, the CO2 molecules were in free motion, leading to an expected disorder. Detailed analysis from the crystal structure indicates that CO2 molecules are predominantly located in the (004) plane. Therefore, the peak position of the (004) plane in PXRD at RT shifts to a higher 2θ when CO2 is released (Figure S5).
The effect of counter ions on formation of CO2 clathrate
A series of guanidinium salts were selected to be tested, including guanidine hydrochloride, guanidine nitrate, and guanidine dihydrogen phosphate. Saturated solutions of guanidinium salts were added into glass vials, which were separately placed in an autoclave. The air in the autoclave was exchanged with CO2 at least three times before the CO2 pressure inside was fixed at 1,000 kPa (10 bar). The autoclave was kept at 0°C, controlled by a cooling jacket. After 24 h with continuous stirring, glass vials were checked for any precipitation.
The pH influence on CO2 adsorption
The pH value of Gua2SO4 aqueous solution is determined to be 7. The pH values of a series of Gua2SO4 solutions (73.7 wt %) were set at 1, 2, 3, 4, 5, and 6 using H2SO4, HCl, and HNO3 solutions. Solutions were then charged into separate glass vials and weighed. The vials were put in an autoclave that was already connected to a CO2 cylinder. The air in the autoclave was exchanged with CO2 at least three times before the CO2 pressure inside was fixed at 1,000 kPa (10 bar). The temperature of the autoclave was kept at 0°C using a cooling jacket. After 12 h of constant stirring, the pressure was recorded, and the vial containing Gua2SO4 aqueous solution was weighed again to calculate the adsorbed amount of CO2. Note that the dissolved CO2 was ignored in calculations, as the CO2 dissolved in solution is negligible (1.44 mg/g at 25°C) compared with that of the amount adsorbed by Gua2SO4.
The effect of pressure on the CO2 uptake of Gua2SO4
A Gua2SO4 aqueous solution (60 wt %, 1.5 g/g) was charged into a glass vial and weighed. The vial was put into autoclave that was connected to a CO2 cylinder. The air in the autoclave was exchanged with CO2 at least three times before the CO2 pressure inside was set to the desired values. The temperature of the autoclave was kept at 0°C using a cooling jacket. After 12 h of continuous stirring, the pressure was recorded, and the vial containing Gua2SO4 aqueous solution was weighed again to calculate the conversion rate. Note that the dissolved CO2 was ignored in the conversion calculations, as the CO2 dissolved in solution is negligible compared with that of the amount adsorbed by Gua2SO4.
CO2 equilibrium pressures of CO2-Gua2SO4 aqueous solutions at various temperatures
A saturated Gua2SO4 solution at 40°C (10 mL) and 2 g Gua2SO4 powder was mixed in a 100 mL autoclave under constant stirring. The temperature of the autoclave was kept at 0°C using a cooling jacket. The autoclave was evacuated under vacuum and filled with CO2. The flushing was repeated three times to completely remove air from the autoclave. Then, the autoclave was charged with CO2 to 100 kPa (1 bar). The pressure drop was monitored until readings were steady. The final pressure value was taken as the equilibrium pressure at the set temperature (0°C, 5°C, 10°C, 15°C, 20°C, 25°C, and 30°C). The water vapor pressure of a saturated Gua2SO4 solution was also measured under the same conditions. CO2 equilibrium pressure was obtained by subtracting the water vapor pressure at the same temperature from the recorded total pressure and plotted against temperature (Figure S11).
Cycling sorption experiments
2.5 g Gua2SO4 was dissolved in a glass vial with 2 mL water. The vial was put in a 100 mL autoclave, and the temperature was kept at 0°C using a jacket-cooling system. The air in the autoclave was exchanged with CO2 at least three times before the CO2 pressure inside was set to 1,000 kPa (10 bar). The pressure was monitored using a digital pressure meter (5 s per point) with constant stirring. After the pressure stabilized, the chamber was vented, and white slurry was observed in the vial. Ultrasound was then applied to decompose the product and release CO2, with a rush of bubbles indicating the decomposition process. The procedure was repeated to test the cyclic performance of CO2 adsorption for at least ten runs. After the final cycle, absolute ethanol was added to the solution, and a precipitate was obtained, which was subsequently dried at 100°C. The PXRD of the precipitate is consistent with that of the pristine Gua2SO4 (Figure S14). The conversion rate was determined gravimetrically by weighing the Gua2SO4 aqueous solution before and after CO2 adsorptions. The conversion rate can also be calculated according to the CO2 pressure drop during the sorption process. Control experiments were conducted following the same procedure but using pure water instead of Gua2SO4 aqueous solution, and N2 gas instead of CO2, and with or without stirring.
Isochoric adsorption-desorption experiments
2.5 g Gua2SO4 was dissolved in a glass vial with 2 mL water. The vial was put in a 100 mL autoclave, and the temperature was kept at 0°C using a jacket-cooling system. The air in the autoclave was exchanged with CO2 at least three times before the CO2 pressure inside was set to 1,000 kPa (10 bar). The pressure was monitored using a digital pressure meter (5 s per point) with constant stirring. After 3 h, the pressure became stable. Then, the temperature of the autoclave was increased to 30°C using a water bath, and the pressure inside the autoclave was recorded.
CO2 removal from flue gas
Owing to the relatively high viscousity of concentrated Gua2SO4 solutions, an autoclave equipped with mechanical stirring was used in this experiment to speed up CO2 dissolution and diffusion. Gua2SO4 (27 g) and water (7 mL) were charged into the autoclave. The air in the autoclave was first exchanged with flue gas three times before setting the gas pressure inside to a specific value. The temperature of the autoclave was kept at 0°C using a cooling jacket. CO2 content after the sorption for 24 h was analyzed by a GC. The removal percentage of flue gas under different pressure is shown in Table S5.
Acknowledgments
We acknowledge support from the Chinese Academy of Sciences, the National Key Research and Development Program of China (2021YFA1500402), the National Natural Science Foundation of China (NSFC; 21571167, 51502282, and 22075266), and the Fundamental Research Funds for the Central Universities (WK2060190053 and WK2060190100).
Author contributions
B.L. conceived of the idea, and B.L., Q.X., and C.T.Y. supervised the project together. Z.X. designed and carried out the experiments. C.L., X.X., and T.S.N. contributed to data analysis and manuscript preparation. C.C. collected SC XRD data and solved the structure. B.L., Q.X., and C.T.Y. wrote the manuscript together. All authors discussed the results and assisted with manuscript preparation.
Declaration of interests
C.T.Y. is an advisory board member at Cell Reports Physical Science. USTC filed one Chinese patent (application #202210271224.0) from the data reported in this study.
Sumber :
https://www.sciencedirect.com/science/article/pii/S2666386423001510
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